SKI Report 96:38

Mineral Formation on Metallic Copper in a "Future Repository Site Environment"

Örjan Amcoff Katalin Holényi

April 1996

ISSN 1104-1374 ISRN SKI-R--96/38--SE

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STATENS KÄRNKRAFTINSPEKTION VOL 2 7 fia 1 5 Swedish Nuclear Power Inspectorate SKI Report 96:38

Mineral Formation on Metallic Copper in a "Future Repository Site Environment"

Örjan Amcoff Katalin Holényi

University of Uppsala Institute of Earth Sciences, Mineralogy-Petrology, Norbyvägen 18B, 752 36 Uppsala, Sweden

April 1996

This report concerns a study which has been conducted for the Swedish Nuclear Power Inspectorate (SKI). The conclusions and viewpoints presented in the report are those of the authors and do not necessarily coincide with those of the SKI. NORSTEDTS TRYCKERI AB Stockholm 1996 SUMMARY.

Formation of copper minerals in a "future repository site environment" is discussed. Since reducing conditions are expected much effort has been concentrated on Cu-sulphides and CuFe-sulphides. However, oxidizing conditions are also discussed. A list of copper minerals are included. It is concluded that mineral formation and mineral transitions on the copper canister surface will be governed by kinetics and metastabilities rather than by stability relations. The sulphides formed are less likely to form a passivating layer, and the rate of sulphide growth will probably be governed by the rate of transport of reacting species to the cansiter surface. A series of tests are recommended, in a milieu resembling the initial repository site conditions.

SAMMANFATTNING.

Bildning av kopparmineral i en framtida slutförvarsmiljö diskuteras. Eftersom reducerande betingelser kan förväntas har Cu-sulfider och CuFe-sulfider getts stor plats, men även oxiderande betingelser diskuteras. En lista på kopparmineral är inkluderad. Slutsatsen är att mineralbildning och mineralförändring på ytan av kopparkapslarna kommer att styras av kinetik och metastabiliteter snarare än av jämviktsrelationer. Det är mindre troligt att sulfiderna kommer att bilda ett passiviserande skikt och sulfideringshastigheten kommer troligtvis att bestämmas av hur fort reagerande ämnen kan transporteras till kapselytan. En serie tester rekommenderas, i en miljö som liknar den initiala slutförvarsmiljön. TABLE OF CONTENTS.

1 INTRODUCTION 1 2 FORMATION OF MINERAL PHASES ON METALLIC COPPER 3 2.1 General 3 2.2 The Cu-Fe-S system 5 2.3 "Problems" concerning low temperature relations in the Cu-Fe-S system 7 2.4 Mechanisms of sulphide formation on metallic copper 13 2.5 "Oxide" formation on metallic copper 1 6 2.6 What phase(s) will form on the canister surface? 20 2.6.1 High sulphur activity conditions 23 2.6.2 Low sulphur activity conditions 24 2.7 Rate of sulphidation on metallic copper 25 3 CONCLUSIONS AND RECOMMENDATIONS 2 8 4 TABLE 1 (copper minerals) 31 5 REFERENCES 37 6 APPENDIX 45 6.1 General 45 6.2 Persistency field Eh-pH relations 49 7 FIGURE CAPTIONS 56 1 INTRODUCTION.

The present report was initiated by the Swedish Nuclear Power Inspectorate. It may be regarded as a continuation of our previous report; "Stability of metallic copper in the near surface environment" (SKN 57, 1992), which dealt with geochemical aspects of copper solubility and stability in an expected future "repository site environment for spent nuclear fuel".

Here we discuss possible mineral forming processes on the surface of the copper canisters, based on the geological - geochemical literature. The advantage in using geological analogies is that nature provides us with long-term low temperature experiments on mineral behaviour, which we can never hope to copy in the laboratory. The disadvantage is that the conditions, in terms of chemical composition etc., are often ill defined.

Unfortunately, information which has a direct bearing on the present well-defined problem is missing, to a large degree. This means that we have to use circumstantial evidence and analogy reasoning, to be able to say anything about possible futures for the canisters. Partly, this problem stems from the fact that studies on ore formation - ore alteration etc., are based on stability relations of various ore minerals at elevated temperatures. Here we are confronted with the problem of predicting what to expect, in an environment (about +5 to +80°C) where most ore minerals fail to equilibrate in the laboratory within reasonable times (This does not mean that the minerals fail to react at all, rather that reactions tend to follow metastable paths at low temperatures. This is true both for natural long-term - and laboratory short- term experiments, and is discussed below).

Thus, studies on mineral equilibria are of limited value, since we have to extrapolate experimental results from, say 200 - 300°C, to the present environment. To add to the uncertainty, it is evident that even equilibrium Eh-pH studies, which are frequently employed to delineate mineral stabilities in the natural low-temperature environment, often show large discrepancies when confronted with reality. As exemplified in Appendix 1, Eh-pH values calculated from thermochemical data, often differ significantly from what is actually observed employing potentiometric measurements.

Since considerable uncertainty exists as to which minerals (if any!) will actually form on the copper canisters, we have listed both likely and less likely candidates in Table 1. Much of the discussion here is concerned with Cu-sulphides and CuFe- sulpjiides. This is partly so because we are more familiar with these phases compared with Cu-oxides and complex Cu(n)-compounds like malachite;

Cu2(CO3)(OH)2 etc., found in environments suggesting oxidizing conditions. Although measurments on deep groundwaters indicate reducing conditions, it is also of interest to analyse the behaviour of copper under oxidizing conditions. Oxidizing conditions will prevail in the repository for some time due to residual oxygen from entrapped air.

Formation and alteration of metal sulphides in nature is a question of redox processes. Some basic definitions and basic thermodynamic quantities were given in SKN 57, 1992, as well as a few examples on how to construct a Eh-pH stability diagram. This is not repeated here. Besides, in the present report we discuss qualitative rather than quantitative aspects of mineral formation on the copper canisters. Quantitative aspects depend on the environment, in terms of exact chemical composition of the groundwater and behaviour of the other barriers (bentonite, enclosing rocks). These questions are outside the scope of this report.

Much of the text below is given to exemplifications and discussion of low temperature processes observed in nature. As stated above, several of these examples are not directly applicable to the present problem. We think they merit our attention, however, since low temperature processes among metal sulphides show some common behaviour, which can be applied to the question of mineral formation on metallic copper. Also, they illustrate the complexity of the processes to be expected, as well as the difficulty confronting us when we try to make predictions about something we have little solid knowledge about. 2 FORMATION OF MINERAL PHASES ON METALLIC COPPER.

2.1 General.

As discussed by Amcoff and Holényi in SKN 57, 1992, there are two major environments in nature where metallic copper is found; (i) In meteorites, "Moon- rocks", and mantle rocks (magmatic rocks with a low SiO2 content, crystallized at great depth), which suggest the strongly reducing environment when the Earth was formed, (ii) In the deeper parts of the oxidized zone of some Cu-rich sulphide ores (still in the upper part of the crust), where conditions are not oxidizing enough for formation of Cu-oxides etc., and not reducing enough for Cu-sulphide formation. It was also suggested that environment (ii) is more relevant when discussing the future repository site environment.

As far as we know, geochemical studies concerning the present problem are practically non-existent. This is so, because the interest has been focused more on the particulars of metallic copper formation, in terms of the physico-chemical milieau, than on its subsequent surface alteration characteristics. The reason for this is partly that metallic copper is a rare mineral, its very existence has been considered the most interesting discussion topic. Besides, microscale surface processes have earlier been difficult to study, they have consequently recieved little attention.

Thus, information which is of interest in the present context is of a rhapsodic and indirect character. As seen in Table 1, there exist many Cu-minerals in the low temperature environment. Some of these may be of importance here.

A number of factors influence mineral formation in the present environment. These include; (i) Temperature and pressure, (ii) Chemical environment in terms of aqueous and solid species which may participate in reactions with the metal, (iii) Stability-metastability. Again we would like to stress that the present report is concerned with qualitative aspects of mineral formation. (For example, the redox conditions in terms of Eh tell us little about the redox capacity, which depends on how fast redox couples which participate in reactions are replenished. This quantity is of fundamental importance when it comes to the question of long-term survival of the copper canisters).

For the sake of argument we assume that the groundwater surrounding the canisters contains a certain amount of species with the ability to oxidize metallic copper; that some sulphur is present in the form of aqueous sulphides and polysulphides (discussed below), and that some dissolved iron is present (SKB 86-03, 87-07). Also we assume a temperature below 100°C and a pH slightly on the alkaline side of neutrality. Here we will not discuss the possibility that the copper canisters might accidentally be subjected to strong oxidation "from above".

A sulphidation process is equivalent with a redox process, independant of whether it takes place in an aqueous or a dry environment. Copper in binary Cu-sulphides seems to be univalent, independent of the Cu/S ratio (Rolf Berger, Institute of Chemistry, Uppsala University, personal communications). This is the case also for CuFe-sulphides, where the valency of Fe varies (+2 - +3), depending on composition, as observed in a number of Mössbauer studies. (for example,

1+ 3+ 1+ 3+ 1+ 2 5+ Cu Fe S2, in chalcopyrite, Cu 5Fe S4, in bornite, and Cu Fe - 2S3, in cubanite. In idaite; approx. Cu5FeS6, which is often found together with CuS, it is

1+ 3+ 4y 3 Tnis likely that the valency of sulphur varies according to; Cu 5Fe S ' "6. phase has not been investigated, however. ). Crystals of Cu-sulphide or CuFe- sulphide, formed on the canister surface, may be subjected to further change through replacement, a process discussed below. In most cases electron transfer (redox reactions) is involved in replacement among metal sulphides, but it may also be a simple ion exchange process.

Before we start to discuss mineral formation mechanisms, however, we shall take a look at the ternary Cu-Fe-S system, since phases belonging to this system are most likely to form on the canister surface. We will concentrate somewhat on the ternary phases (minerals containing both Cu and Fe), since aqueous Fe-species are important in most natural environments. 2.2 The Cu-Fe-S system.

This system has probably recieved more attention than any other ternary metal sulphide system (e.g. Merwin and Lombard, 1937; Brett, 1964; Yund and Kullerud, 1966; Cabri, 1973; Sugaki et al, 1975; Vaughan and; Craig, 1978; Durazzo and Taylor, 1982; Kojima and Sugaki, 1987). This is so because CuFe-sulphides are

almost invariably present in sulphide ores (chalcopyrite; CuFeS2, is one of the most widespread metal sulphides on Earth), and because they show a great deal of complexity and interesting behaviour, especially at low temperatures.

Since our knowledge of the complicated low temperature relations in the system are based on extrapolations from experiments performed at elevated temperatures, and on studies of natural occurences, we have chosen to discuss the relations at decreasing temperatures (Figure 1). High temperature relations have no direct bearing on the present problem, but they are valuable to know anyway, because they indicate the direction towards low temperature relations (Figure 2), which are considerably more uncertain because of problems with metastabilities etc.

Our knowledge of these phases are mainly based on sc called "dry experiments", i.e. pure metals and elemental sulphur in desired proportions are reacted at high temperatures in vacuum until equilibrium is achieved. Comparative studies where water is included, using pressure vessels, give similar results, at least at high temperatures. Metal sulphides are much more sensitive to temperature changes than to pressure changes. For those unfamiliar with phase diagrams, it should be noted that the lines between different phases in Figures 1 and 2 (tie-lines), connect phases which are supposed to be in thermodynamic equilibrium with each other under the specified conditions. (The so called thermodynamic phase rule states that a maximum of five phases can co-exist stably with one another in a ternary system, given that both temperature and pressure are exactly fixed. In the real world, however, both these variables vary, which gives a maximum of three co-existing phases, as seen in Figures 1 and 2).

Two things are apparent when the phase relations at different temperatures are

compared; (i) The complexity increases when temperature is decreased, and (ii) Extensive solid solutions are found at high temperatures, while only phases with relatively well-defined compositions are stable at room temperature in the present system. This is due to the fact that Cu and Fe atoms are randomly distributed in tetrahedral intersticies or holes in a close-packed sulphur structural framework at high temperature (Figure 3). In a close-packed structure the number of tetrahedral intersticies is twice the number of atoms making up the close-packed framework. These factors make it possible to vary both the Cu/Fe ratio and the metal/S ratio within certain limits, depending on the specific structure. At low temperature, on the other hand, Cu and Fe tend to order with respect to one another, for energetic reasons, resulting in a number of possible structures with distinct Cu/Fe - and metal/S ratios {superstructures).

At 600°C the system is dominated by three extensive solid solution fields; (i) a pyrrhotite; Fe-|.xS, series (poss), which can accomodate several wt.% Cu, (ii) an intermediate solid solution series (iss) which extends compositionally from about

CuFeS2 to CuFe2S3, on the metal-rich side of a 1:1 metal:sulphur ratio, and (iii) a bornite solid solution series (bnss), which extends from the Cu-S binary, at about

Cu2S - Cui 7S, to ternary compositions half-way towards CuFeS2 (Figure 1). The only other phase which is stable at this temperature, besides the elements themselves, is pyrite; FeS2.

As temperature is decreased, chalcopyrite; CuFeS2, appears at 557°C (a superstructure of iss), covellite; CuS, at 507°C, and idaite; Cu5FeS6, at 501°C.

The solid solution fields decrease successively in size, and bnss is split into two solid solutions at around 300°C, one Cu-rich; digenite solid solution (dgss) and one Fe- rich (bnss) (Craig and Scott, 1974).

Between 300°C and room temperature a large number of changes take place. The solid solutions, where the metal atoms are disordered, break down completely, and a large number of superstructures appear. For example, within the compositional field occupied by the cubic iss at high temperature, we find at least five phases with well-defined compositions; chalcopyrite; CuFeS2, cubanite; CuFe2S3, talnakhite;

CugFegS-|6, moohekite; CugFegS-jg, and haycookite; Cu8FeinS16. Tetragonal bornite; Cu5FeS4, appears in place of the cubic bnss, and dgss breaks down to chalcocite; Cu2S and digenite; about Cu-| 8S (possibly stabilized by a minor amount of Fe) (Figure 2). Additional phases appear along the Cu-S binary, and several varieties of covellite (CuS) have been found in nature. Also, nunkundamite; about Cu-|iFe2S-|3, may form in stead of idaite; Cu5FeS6 (Wang, 1984). (Note that both these phases are "covellite-like" in appearence).

2.3 "Problems" concerning low temperature relations in the Cu-Fe-S system.

Since we are here interested in what will actually happen at the metallic copper - hydrous environment interface, it seems appropriate to expand the discussion a bit, regarding low temperature mineral formation. We are confronted with large discrepancies between on the one hand laboratory results and on the other hand what we observe in nature. Below, we will give a few examples, illustrating that low temperature phases found in nature may have formed and been altered along quite different reaction paths, as compared with phases produced in the laboratory. Not all of these examples are directly relevant in the copper canister context, but we believe they are instructive in showing the problems involved in making low temperature predictions.

(1) Cubanite; CuFe2S3, is an orthorombic mineral found in many copper ores, often intimately (=crystals intergrown on the microscale) associated with chalcopyrite, and pyrrhotite. It is stable to 200-210°C, above which it rapidly transforms to cubic isocubanite (with Cu and Fe disordered), which is a member of the high temperature intermediate solid solution (iss). Although reaction rates are relatively rapid at 200°C, the cubanite -> isocubanite transition has never been successfully reversed in the laboratory. In fact, it has not been possible to synthesize cubanite under any laboratory conditions.

According to Putnis and McConnell (1980) this enigmatic behaviour can be explained as follows: In the cubanite structure sulphur exhibits hexagonal close- packing, changing to cubic close-packing (Vaughan and Craig, 1978) in isocubanite. A change in sulphur packing is equivalent with a major transition and sulphur behaves as a relatively sluggish component in all kinds of processes taking place in metal sulphides. It is possible to bring about the cubanite -> isocubanite transition 8 since we can use excessive heat, but it is not possible to bring about its reversal since we must keep the temperature below 200°C. The interesting point is that isocubanite decreases its free energy on cooling by exsolving chalcopyrite; CuFeS2, from the cubic matrix. In chalcopyrite sulphur exhibits cubic close-packing, and is well-ordered at variance with isocubanite. The isocubanite "motherphase" becomes correspondingly enriched in Fe during the exsolution process. This is typical metastable behaviour, and the decrease in energy which accompanies the exsolution process further decreases the tendency for isocubanite to transform to cubanite. (Exsolution in the solid state = precipitation of a supersaturated solid phase from a solid solution. Normally it takes place on cooling.).

Cubanite (CuFe2S3) in nature is found in environments suggesting strongly reducing conditions (see type a environment, page 4 in SKN 57, 1992) and a rock-forming temperature much in excess of 200°C. Textural relations (= how the mineral grains are intergrown with each other) between cubanite (CuFe2S3), chalcopyrite

(CuFeS2) and pyrrhotite (Fe-|.xS), suggest that cubanite has formed at quite a low temperature, however, by means of chemical replacement (A process where compositional and structural changes take place in a solid phase at practically constant volume. Examples are given in Figure 4.).

We can consider a hypothetic redox process:

2+ 1 + 3CuFeS2(ox) + Fe + e* -> 2CuFe2S3(red) + Cu

1+ 2+ - 3FeS(red) + Cu -> CuFe2S3(ox) + Fe + e which gives the overall process:

CuFeS2 + FeS -» CuFe2S3

Among metal sulphides, such replacement processes, which are very complex in detail, take place by means of rapid electron transfer together with diffusion of metal ions through an essentially static sulphur structural framework. The reaction rate 9 is probably increased by presence of water.

(2) Bornite; Cu5FeS4, is a tetragonal phase, which exhibits similar interesting behaviour. On heating to 228°C (Craig and Scott, 1974) it inverts to a cubic disordered form, which is a member of bnss. Rapid cooling of cubic bornite leads to formation of cubic superstructures (ordering takes place), whose unit cell dimensions are simple integers of the unit cell of the high temperature form, {unit cell = the smallest possible repetitive building block of a crystal. All information necessary to define the crystal structure is contained in the unit cell. For a cubic crystal, for example, it constitutes a small cube). However, at variance with cubanite, the stable tetragonal form of bornite forms on slow cooling below the transition temperature. This is so because sulphur exhibits cubic close-packing in both crystal structures, i.e. no major structural reorganization is needed to bring about the change, only ordering between Cu and Fe.

In nature, bornite is found in primary copper ores formed at low temperature in the near surface environment, as well as in ores which have been subjected to secondary alteration (=redox processes essentially, which take place because of interaction with groundwaters). Both environments can be considered to correspond with environment b, page 5 in our previous report; SKN 57, 1992. Bornite is often found together with secondary Cu-sulphides (Cu2S, CuS etc), and chalcopyrite and digenite are commonly present as beautiful, crystallographically oriented two- phase "exsolution lamellae" in these bornites. (During exsolution the "daughter" phase commonly grows only along certain energetically favourable crystal planes in the "mother" crystal, resulting in intricate network textures. Most prominent in this respect is the so called Widmannstätten texture in iron meteorites (Ramdohr, 1980) where two FeNi alloys have formed on slow cooling of a single high temperature solid solution phase). Moreover, the lamellae are typically rimmed by digenite; Cu1 8S (Figure 4).

Indeed, cooling of experimentally produced crystals of bornite solid solution; bnss,

enriched in CuFeS2 compared with a nominal Cu5FeS4 composition, invariably results in lamellar exsolution of chalcopyrite from the "mother" bornite phase.

However, at temperatures in excess of about 200°C crystallographic coherency is 10 lost in a very short time, and the chalcopyrite lamellae recrystallize to rounded grains in the bornite host crystal (e.g. Durazzo and Taylor, 1982).

Thus, the "exsolution lamellar" texture observed in nature point to near room temperature conditions. Besides, bulk chemical analyses of bornite crystals, including exsolved phases, points to compositions well outside the boundary of the high-temperature bornite solid solution; bnss. In analogy with formation of cubanite

(CuFe2S3), this indicates that these intergrown bornite - chalcopyrite - digenite

(Cu5FeS4 - CuFeS2 - Cu-j 8S) crystals have formed by means of chemical replacement, i.e. replacement exsolution. We can consider a simplified reaction:

1+ Cu5FeS4(red) -> CuFeS2+2Cu1 8S(ox) + 0.4Cu + 0.4 e" taking place during slight oxidation at near room temperature conditions (Amcoff and Hedin, 1992). Dissolved copper ions precipitate as CuS along grain boundaries, while chalcopyrite and digenite form lamellae in the host bornite crystals.

Microscopic observations (e.g. Ramdohr, 1980) and stability considerations, suggest that chemical reactions among sulphide minerals at low temperatures to a large degree is a question of chemical replacement. For Cu- and CuFe-sulphides this means a redox process where metal atoms and electrons are rapidly exchanged through an essentially static sulphur structural framework, i.e. the reaction takes place with a minimum of structural reorganization (Amcoff, 1987; 1988; Amcoff and Hedin, 1992).

(3) Stability relations among Cu- and CuFe-sulphides at low temperature are uncertain. For example, Craig and Scott (1974) proposed that the common

association of pyrite and bornite (FeS2 + Cu5FeS4) at Magma Mine, Arizona, at

present rock temperature of 50°C (a long-term natural equilibrium experiment!), suggests that the two phases may co-exist stably at room temperature. In Sweden, the two minerals are found in direct grain contact, for example at Mårtansberg (Geijer, 1924). 11

On the other hand, Yund and Kullerud (1966) proposed that a tie-line; bornite - pyrite (Cu5FeS4 - FeS2) is replaced by a tie-line; chalcopyrite - digenite (CuFeS2

- Cu-j 8S), below 228°C. This was established experimentally, both by reaction; pyrite + bornite to produce chalcopyrite + digenite, slightly below 228°C, and by exsolution of the same product phases from sulphur enriched bornites at low temperature.

Moreover, hydrothermal experiments by Sugaki et al (1975), using pressure vessels, indicated a tie-line; idaite - chalcopyrite (Cu5FeS6 - CuFeS2), at 300 -

350°C, but Kojima and Sugaki (1985) showed that a tie-line bornite-pyrite

(Cu5FeS4 - FeS2) exists in the Cu-Fe-Zn-S system, at least above 300°C. (Note that Zn does not enter the Cu-minerals to a large extent. Also, at these temperatures equilibrium should be achieved relatively rapidly).

Additional examples showing discrepancies among Cu-mineral stability relations could be given. These findings all suggest small differences in energy between different alternatives, and metastable behaviour is to be expected. Also, different results between anhydrous and hydrothermal experiments, suggest that H-S bonds in sulphides may play a role. Besides, the different observed results between the Cu- Fe-S and Cu-Fe-Zn-S systems (see above) suggest that trace amounts of Zn in the Cu-mineral structures may have a strong influence on the stability relations. It may be added here that a large number of elements are invariably present in trace amounts in naturally occuring metal sulphides. In fact, pure phases are absent in nature for energetic reasons.

(4) A number of metal sulphides found in nature appear to be impossible to synthesize in the "dry" system, for reasons other than discussed in connection with the "cubanite enigma" above. Several such Fe-sulphides are found in the "low temperature sedimentary environment" (stagnant basins on the seafloor, suggesting reducing conditions). Mackinawite (FeS with tetragonal PbO structure), greigite

(Fe3S4; a thiospinel), and marcasite (orthorombic metastable FeS2; compare with pyrite which is cubic FeS2) are found in this milieau, as well as additional more or less well-defined phases (e.g. Berner, 1964; Rickard, 1975). These Fe-sulphides 12 can be synthesized in a water system of appropriate chemical composition at low temperature. According to Rickard (1969, cited in Rickard, 1975) marcasite tends to form under acid conditons and pyrite under more alkaline conditions.

According to Kullerud (1967, cited in Craig and Scott, 1974) marcasite can be inverted to pyrite down to 150°C in the presence of excess sulphur, but only above

400°C in the absence of sulphur; this suggests a slightly higher Fe/S ratio in marcasite. From his experiments Kullerud (1967) proposed that H-S bonds might stabilize the marcasite structure. (It is impossible to convert pyrite to marcasite in a "dry" experiment at any temperature.).

To add to the complexity, a large number of pyrrhotite varieties of compositions between troilite (stoichiometric FeS) and monoclinic pyrrhotite (Fe7S8) have been reported in the literature, all of which may be metastable with respect to the end- members, according to Putnis and McConnell (1980). The conclusion to be drawn here is, that there probably exist many more metastable Fe-sulphides in nature than the thermodynamically stable ones.

To conclude: The reason for discussing the "natural Fe-S system" here is that it is the most studied among binary sulphide systems (Fe-sulphides are most important, quantitativly). There is no reason that Cu- or CuFe- sulphides should be more simple at low temperatures. In fact, reaction rates generally increase when the Cu- content increases among phases in the Cu-Fe-S system, indicating that relations could possibly be even more complex along the Cu-S join compared with the Fe-S join. Textural interrelations (=how the mineral grains are intergrown) among bornite (Cu5FeS4), chalcocite (Cu2S), digenite (Cu1 8S) etc., are extremely diverse, suggesting a variety of possible processes, depending on the conditions. Colour variations in Cu-sulphides observed under the microscope, point to minor amounts of eg. Ag, Fe, Mn in the solid matrix. It may be noted here that in mineral samples subjected to air, tarnishing sets in immediately. The compositions of the thin films formed are little known, but Ramdohr (1980) suggested that a thin layer of Cu-j 8(S,O) forms on Cu2S crystals subjected to air. Instant tarnishing on synthetic crystals of bornite and chalcopyrite has also been observed by one of the present authors, indicating transport of Cu to the surface of the grain and 13

precipitation of oxidized phases (e.g. CuS, Cu5FeS6, and possibly Cu-oxides). Once started, tarnishing continues for a while, even when the sample is placed under acetone!

2.4 Mechanisms of sulphide formation on metallic copper.

Whatever the subsequent history, an electron transfer from metallic copper must take place in the initial sulphidation step. Copper is oxidized from zerovalent to univalent. A suitable electron acceptor must be present at the canister surface. A number of possible alternatives are present in various groundwaters, which we will not dwell on here, but it may be noted that even aqueous sulphide species may oxidize metallic copper. Thus, aqueous polysulphide ions, the formation of which can be

2 considered as a reaction between S°(aq) and aqueous H2S, HS" or S ' (depending on pH), have mean valencies between 0 and -2 (Shea and Helz, 1988), and will most likely be present in the repository environment (Nordstrom SKB 86-03). According to Nordstrom, a polysulphide zone will probably be present above a sulphide zone in a typical Swedish "granitic groundwater", reflecting the somewhat higher oxidation state. From a mechanistic point of view (kinetic collision theory), it seems feasable that a polysulphide ion colliding with the canister surface, might act both as an electron acceptor and as a source of sulphur. Formation of FeS2 from FeS by a similar process was indicated by Rickard,1975; discussed below). As the polysulphide/"simple sulphide" ratio decreases, other oxidants in the neighbourhood may oxidize "simple sulphide" to polysulphide, to maintain the reaction. Besides,

also in the presence of "simple" sulphide species (h^S, HS") Cu-sulphides may

+ form, since H will act as an electron acceptor leading to H2(g) formation.

However, once a solid sulphide "layer" has precipitated on the metal surface, continuing reaction will take place by electron transfer and copper diffusion through the semiconducting meta! sulphide to the surface, while aqueous sulphur species as well as other possible reactants, like ferrous iron, will diffuse through the surrounding aqueous medium to the reaction site (Figure 5). This will continue as long as reactants are available in the neighbourhood. 14

We do not think it is possible to show the redox process in detail at present. As already hinted at, there will probably be several alternatives among the aqueous species with the capacity to oxidize metallic copper. Furthermore, as discussed in Appendix 1, redox processes among aqueous species are often path controlled, leading to formation of metastable products. Processes involving transfer of many electrons (more than one or two) are disfavoured as are also processes where strong chemical

2 bonds are broken. For example, the thiosulphate ion; S2O3 ", acts as a metastable

2 2 but persistent intermediary, in both oxidation and reduction, for the S 7SO4 " redox couple (Ohmoto and Lasaga, 1982; Sato, 1992). Also, hydrogen peroxide;

H2O2, acts as an intermediary for reactions involving the important O2/H2O redox couple (Moffett and Zika, 1983; Millero et al, 1987; Sato, 1992), which defines the upper Eh boundary for the natural environment.

Regarding crystalline phases at low temperature, the discussed replacement examples above demonstrate that chemical change takes place in such a way that a minimum of structural reorganization is involved. In metal sulphides the metal atoms are typically very mobile, while sulphur acts as a sluggish component. Chemical replacement often results in formation of metastable phases and textures.

Thus, sulphidation of metallic copper in a repository site environment is a process which involves an outward diffusion of copper and precipitation on the surface of the Cu-sulphide crystal, growing into the bentonite surrounding. The details of the continuing growth process are not obvious but we think dendritic growth through the bentonite layer is possible {dendrites are single crystals which form during growth from a non-convective supersaturated solution, which can be a melt, a solid phase or a gas phase). As a crystallite grows by means of diffusion of solute species through the solution, the neighbourhood to the crystal will be depleted in these species and concentration gradients are formed. Since a crystal face will need more solute particles per unit area for a certain amount of growth, compared with a crystal edge, which in turn will need more solute particles compared with a crystal corner, the concentration gradient around the crystal will become steeper in the order; crystal face < crystal edge < crystal corner. This will favour growth in the order; crystal corners > crystal edges > crystal faces, which may result in dendritic growth. In a saturated solution dendrites will tend to change into normal 15

crystals since the precipitating particles are more strongly bonded to a crystal face than to a crystal corner. Also, for the same reason, dissolution in a non-convective solution will tend to operate in such a way that corners are dissolved before faces, leading to rounded crystals. Strong undersaturation may lead to dissolution from dislocations etc. on the crystal faces, resulting in the appearence of etch pits. Since bentonite constitutes an unconsolidated mass of clay mineral grains which are only loosely connected with each other, Cu-sulphide crystals will probably grow into the space between individual clay particles and encapsulate them. This is a common process during crystal growth in nature; large single crystals filled with inclusions of other "engulfed" phases are called poikiloblasts. However, in this way we think growth of Cu-sulphides in a "wet" bentonite medium may resemble open space growth rather than arrested growth in a confined space. Experimental sulphidation by Amcoff (1987, 1988) in "open space" showed that covellite and idaite crystals, forming on the surface of the phase which was sulphidized, typically developed as single crystals which always grew outward in a direction perpendicular to the surface from where reaction started. Considering such a "single crystal" texture as well as the high solid diffusion rates involved, it does not seem likely that the forming Cu-sulphide will passivate the metal.

Cu-sulphides as well as Ag-sulphides react extremily rapid (Barton, 1970), suggesting very high diffusion rates of the metal through the solid matrix. The sulphidation experiments by Amcoff (1987, 1988) also showed, that the position of the original crystal surface (where reaction started) could still be inferred as reaction continued. The process is assymetric in a sense, since the phase which is sulphidized loses copper continuously, while sulphur does not diffuse the opposite way but reacts at the surface of growing Cu-sulphide crystals. Copper is lost from the interior and reprecipitated at the exterior. In this respect it does not seem likely that sulphidation in a water environment will differ much, compared with the discussed dry experiments.

Microscopic examination of natural low temperature Cu- and CuFe-sulphides shows that chemical alteration preferentially proceeds via chemical replacement. In this way "old" crystal shapes may survive, although the chemical composition as well as the crystal structure are changed. Chemical replacement was a popular branch of 16 experimental sulphide mineralogy in the twenties and thirties, and it was repeatedly demonstrated that mm-sized crystals of CuFe-rich sulphides, could be altered within hours into more Fe-rich or more Cu-rich varieties, even below 100°C, by subjecting the crystals to more or less concentrated solutions of suitable copper- and iron salts (e.g. Ray, 1930; Schouten, 1934).

In the copper canister case, replacement will probably become important if and when the chemical composition of the water surrounding the canisters starts to change. For example, the entrapped groundwater may become increasingly depleted in aqueous sulphide species // initial Cu-sulphide precipitation is rapid compared to diffusion through the bentonite barrier to the reaction site (a small redox capacity because of slow replenishment). Thus, an initial Cu-sulphide with a low Cu/S ratio may be replaced by a sulphide with a higher ratio. With time a steady state situation will be set up where diffusion through the solution medium matches precipitation. Another possibility is that a metastable solid phase forms initially, for kinetical reasons, which is subsequently transformed into a thermodynamically more stable mineral by means of replacement.

2.5 "Oxide" formation on metallic copper.

Although the present paper is mainly concerned with sulphide phases associated with metallic copper, formation of other Cu-phases can not be excluded. Cu-oxides like cuprite (Cu2O) and tenorite (CuO), as well as a large number of other complex

CU(II) phases, for example Cu-carbonates like malachite; Cu2(CO3)(OH)2 etc. (see Table 1C, Figures 6 and 7), are found associated with metallic copper in oxidized copper ores. This natural "near room temperature" environment was discussed in our previous report; SKN 57, 1992, pp 5-6. Besides, a number of these phases are formed when copper metal is directly exposed to the atmosphere.

When oxidized copper ore is studied under the microscope (Ramdohr, 1980) a number of conclusions can be drawn which are of interest in the present context:

(i) Metallic copper is often associated with chalcocite (Cu2S) and/or cuprite

(Cu2O). It may replace these phases or be replaced by them, (ii) Replacement involving metallic copper; for example; Cu -> Cu2O or Cu2O -> Cu, is typically 17 massive in character rather than oriented (Figure 4). Thus, grains of metallic copper are often rimmed by cuprite (Cu2O). (iii) Other oxidized phases, like tenorite (CuO), Cu-carbonates and chrysocolla (CuSiO3 + nH2O), as well as limonite; FeO(OH), often precipitate in concentric shells around metallic copper-

Cu2O aggregates, or occur as vein fillings. Banded oxidized copper ore is sometimes quite "massive"; malachite, Cu2(CO3)(OH)2, from the Ural is famous for its use as ornamental stone. Sometimes Cu-carbonates directly replace metallic copper (Ramdohr, 1980) suggesting steep chemical gradients.

The present observations (abundant replacement, rhytmic precipitations, vein fillings, chemical disequilibrium already on the microscale, etc.) suggest a milieu of

+ dissolution and reprecipitation. Cu^ is leached from primary copper ore (CuFeS2) when it is subjected to strongly acid, oxidizing solutions "from above". As sulphuric acid is neutralized by neighbouring silicates and carbonates, a whole spectrum of secondary Cu-minerals, including metallic copper, precipitate and replace each other.

Regarding the present problem it is interesting to note that metallic copper is replaced by Cu2O as well as by Cu2(CO3)(OH)2 in a massive manner. A process is suggested where Cu2O "invades" the metal, i.e. "inward" diffusion of oxygen and

"outward" diffusion of copper through the growing Cu2O layer. A considerable net loss of copper is indicated, since, on a volume for volume basis, Cu2O contains about 40% less Cu than metallic copper. Released Cu probably participates in formation of additional Cu2O in the vicinity. At variance with oxidation replacement of bornite, discussed above, where bornite was replaced by a CuFeS2 + Cu-| 8§ assemblage, a net "inward" diffusion of the anion is indicated in the present case. It is not quite obvious whether this difference reflects a fundamental difference in behaviour between oxygen and sulphur as regards solid state redox processes, or whether it simply reflects the fact that metallic copper is only involved in the Cu -> Cu2O replacement example. To get additional information we will take a look at another example involving oxides:

Since it is very important both from a geological and a technical point of view,

oxidation of magnetite (Fe3O4) has been studied by numerous authors e.g . Davis et 18 al. (1968); Colombo et al. (1868); Feitknecht and Gallagher (1970). The reaction shows some common behaviour with the sulphidation processes discussed above.

The overall redox process can be written:

2Fe3O4 (magnetite) + 1/2O2 -> 3Fe2O3 (hematite)

At low temperature oxidation takes place via intermediate formation of metastable maghemite (Y-Fe2O3; magnetite or spinel structure with hematite composition. A portion of the iron atoms are missing compared with normal magnetite), or of a magnetite - maghemite solid solution (Fe3O4 - Y-Fe2O3):

Fe3O4 (magnetite)-» Fe3_xC>4 - y-Fe2O3 (solid solution) -> Fe2O3 (hematite).

From a mechanistic point of view, this oxidation process takes place via outward

1 diffusion of Fe* "*" ions to form Fe2O3 (maghemite or hematite) on the surface. With time hematite (Fe2O3) also nucleates along (111)-planes in the metastable magnetite-maghemite host crystal, resulting in a typical martitic intergrowth texture (Figure 4). The oxidation replacement process can be written:

2+ 3Fe3O4 -» 4Fe2O3 + Fe + 2e" illustrating that oxygen behaves as an immobile component within the crystal. Note the similarity with oxidation replacement of bornite discussed above:

1+ Cu5FeS4 -> CuFeS2 + 2Cu1 8S + 0.4Cu + 0.4e"

Martitization is a very common process during oxidation of iron oxide ores (for example Kiruna), and is also a characteristic feature during oxidation of Fe -Ti -

oxide ore where ilmenite (FeTiO3) and hematite (Fe2O3) precipitate in lamellar

form along (111)-planes in the host Fe3O4 - Fe2Ti04 (magnetite - ulvospinel solid solution) crystal (e.g. Haggerty, 1976a and 1976b; Lindsley, 1976; Reynolds, 1985 a and 1985b; Amcoff and Figueiredo, 1990). 19

The important point is that the martitization process which is thus an example of

oxidation replacement very much resembles oxidation of bornite (Figure 4).

Additional examples from the Cu-Fe-S system could be given. Thus, no fundamental difference in behaviour between oxides and sulphides is suggested. It is possible that the diffusion rate of copper in a "dense" Cu2O crystal is lower compared with a Cu2S crystal, but we do not have actual figures to corroborate such a statement. (The

martitization process is a more sluggish process compared with oxidation of CuFe- sulphides). Thus, we do not know whether Cu2O will passivate metallic copper.

Initial oxide formation: There is a possibility that Cu-oxide(s) may have formed on the canister surface before sulphidation starts. During construction of the repository it is inevitable that some oxygen will be entrapped, and also a thin oxide coating is likely to have formed already during manufacturing of the canisters. Here it might be added that oxygen and sulphide species are mutually exclusive in a hydrous environment. According to Berner (1981) they react within minutes at

room temperature. (Recently; 5/6 1993, the "dead" H2S-rich bottomwaters of

Gotlandsdjupet, due east of Gotland, were rapidly oxygenated because of influx of O2-

rich heavy "Northseawater", 300 km3 of which was "spilled" into the Baltic during

the strong winterstorms. Since then the O2 concentration has gradually decreased

again, probably because of release of accumulated H2S from the bottom sediments).

A fresh copper surface typically becomes reddish (Cu2O?) in a day or so if subjected to the atmosphere. According to Hägg (1966) a slight heating of copper in

pure air leads to oxidation to Cu2O (yellow-red) and then CuO (black). Subsequently, other compounds may form if the air contains species like sulphuric acid and chlorine compounds etc. Hägg (1966) observed that a copper roof first becomes black from Cu-sulphides. McNeal and Little (1990) observed that Cu-

sulphides (the whole spectrum from Cu2S to CuS) may form on a Cu2O layer on copper metal in an oxidizing environment, because of anaerobic, sulphate-reducing bacteria, forming a surface film. These bacteria are diverse and can be isolated from many aerobic environments, but their principal habitats are seawater and soil. According to Hägg (1966), subsequent processes involving oxidation, result in

formation of greenish Cu(ll)-compounds; especially Cu4(OH)6SO4, but also 20

malachite; Cu2(OH)2CO3 etc. Near the coast, Cu2(OH)3CI may become important. From a mechanistic point of view, this reaction sequence seems logical, since formation of Cu2O or Cu2-S is a relatively straightforward redox process, while formation of e.g. Cu2(CO3)(OH) is a complex process involving many individual reaction steps.

However, an initial thin coating of Cu2O or CuO on the copper canister surfaces will affect subsequent sulphidation at the repository site. (Unless the repository is accidently subjected to strong oxidation for a long period of time, a case which we have not considered here, we do not believe complex Cu(ll)-compounds will form).

As discussed above, it is not quite clear whether the presence of Cu2O could inhibit sulphidation to any significant degree. If Cu2O behaves like a persistent phase, in the repository site milieau, it seems likely that initial reaction with sulphur on the

+ Cu2O crystal surface will be followed by an outward diffusion of Cu through the crystal to the reaction site on the surface, a process not much different from diffusion of Cu+ through a Cu-sulphide. (The rate may differ, however). If, on the other hand, Cu2O does not persist in a sulphidic water environment, it will simply be replaced by Cu-sulphide, i.e:

Cu2O + H2S -> Cu2S + H2O

Such a process is probably quite feasible, although it may be characterized by outward diffusion of copper ions and inward diffusion of protons rather than diffusion of anions through the oxide crystal. Even phases persisting well outside their stability fields are subjected to simple exchange processes (see Appendix 1). Observations from nature have shown sulphidation replacement to be extremely common. According to Bastin et al (1931); " Sulphides replace aN rock minerals, aH gangue minerals (vein fillings etc.), and all sulphides." This shows that these processes are mechanistically favoured.

2.6 What phase(s) will form on the canister surface?

Since we can expect metastabie behaviour it is neither possible to answer in detail 21 which redox processes will be in operation nor which phase(s) will form. Besides, we do not know the exact composition of the groundwater, nor of the enclosing bedrocks. However, in general terms we can relate low temperature metal sulphide behaviour, observed in the experimental Cu-Fe-S system and in nature, to the present problem.

In minor or trace amounts, metal sulphides are almost always present in crystalline bedrocks. Pyrite and pyrrhotite, the two most important iron sulphides, strongly dominate. In sedimentary rocks and in loose sediments, forming at low temperature in reducing environments, other Fe-sulphides like mackinawite; FeS-j.x, may be present as well (e.g. Rickard, 1975). Also, in the bentonite clay a certain amount of iron sulphides may be present. There is a general tendency that more oxidized iron sulphides (FeS2 pyrite) are common in granitic rocks (high SiC^ content), while more reduced iron sulphides (Fe-|_xS, pyrrhotites) dominate in gabbroic rocks (low SiC>2 content). Both types are common in Swedish bedrocks. In magmatic rocks, low in SiO2 (ultramafic rocks), even troilite (stoichiometric FeS) is found. (This is the meteoritic environment discussed above.). These differences reflect increasingly oxidizing conditions in the order; ultramafic rocks -> gabbroic rocks -» granitic rocks.

A considerable amount of data have been collected on groundwater chemistry. For the sake of argument, however, we assume that the composition of the groundwater to which the copper canisters will be subjected, in some way will reflect the composition of the neighbouring bedrock (influx of water which has reacted with sediments from a larger distance, organics and microbial activity affecting the sulphate : sulphide ratio etc., may, of course, also be important). This does not neccessariiy mean that the groundwater has equilibrated with respect to iron sulphides in the neighbourhood. However, we will discuss initial sulphide formation on the canisters in terms of an "iron sulphidic water environment".

Since we do not know whether the environment will be "pyritic" or "pyrrhotitic" or a mixture of both we will discuss both cases. (From Figure 2 it can be seen that neither FeS2 nor Fe-|.xS are in thermodynamic equilibrium with metallic copper at room temperature. On the other hand, troilite; FeS, may probably co-exist with 22 metallic copper at low temperatures. Troilite is not expected in the present environment, however).

(1) Metallic copper subjected to a "pyritic environment": In this case we think the chemical composition of the groundwater will be characterized by a relatively low dissolved iron : aqueous sulphide ratio. From Figure 2a it is apparent that FeS2 may coexist thermodynamically with different Cu-rich sulphide assemblages, i.e. with decreasing sulphur activity; (i) CuS + Cu5FeS6, (ii) Cu5FeS6 + Cu1 8S, (iii)

Cu-j 8S + Cu5FeS4, and (iiii) Cu5FeS4 + CuFeS2. A further decrease in sulphur activity would destabilize pyrite. Note that Figure 2b would yield different assemblages.

(2) Metallic copper subjected to a "pyrrhotitic environment": In this case we think the chemical composition of the groundwater will be characterized by a higher dissolved iron: aqueous sulphide ratio. From Figure 2a and b it is appearant that

Fe1.xS may only co-exist thermodynamically with Fe-rich CuFe-sulphides. With decreasing sulphur activity we have the alternatives: (i) FeS2 + CuFeS2, (ii)

CuFe2S3 + CuFeS2, (iii) CuFe2S3 +FeS. Thus, pyrrhotite is not thermodynamically stable in the presence of simple Cu-sulphides at room temperature. A further decrease in sulphur activity would destabilize pyrrhotite

We have given these examples just to illustrate the complexity of the present problem. It should also be noted that we have here assumed thermodynamic equilibrium. In the real case we must rather assume path-dependant processes, leading to formation of metastable phase assemblages.

Metallic copper will be present in large excess, compared with reactive aqueous species. This means that the groundwater will be depleted in aqueous species, a situation which is quite the opposite the pyritic and pyrrhotitic environments

above, where we assumed aqueous species (resulting from interactions with FeS2

and Fe-j.xS, respectively) in excess. Moreover, if we think of the sulphidation process in terms of kinetic theory (collision rate of reactive aqueous species with

the canister surface), it seems likely that chalcocite; Cu2S, will be the first phase to form, independently of the exact composition of the water. Copper is already present 23 at the reaction site, and a lower collision frequency of sulphide species is neccessary for Cu2S formation, compared with S-richer Cu-sulphides. Also, initial formation of a CuFe-sulphide (e.g. Cu5FeS4) is unlikely, because both Fe and S have to be present simultaneously at the reaction site. This is equivalent with a higher activation energy, or a more sluggish reaction.

The subsequent pathway is more difficult to predict. We will discuss two cases; high and low sulphur activity conditions, respectively. (This thermodynamic quantity depends on the total concentration of available sulphide species).

2.6.1 High sulphur activity conditions.

In this case Cu2S will be transformed into S-richer Cu-sulphides on the surface of the growing crystals. Also, if the groundwater is not depleted in sulphide species, it seems likely that a steady-state process will be set up, where originally formed

Cu2S will survive at the "old" metallic copper - water interface, a whole spectrum of Cu-sulphides will form across the growing crystals, and CuS might be present at the sulphide crystal - water interface (Figure 5). As long as sulphidation continues, we can expect solid state diffusion of copper to determine the phase apparence of the "sulphide layer", where the neighbourhood to metallic copper determines the

"stability" of Cu2S, while the composition of the surface Cu-sulphide is controlled by the composition of the water phase.

This state of affairs becomes obvious when CuFeS2 crystals are experimentally subjected to a sulphur atmosphere (anhydrous experiment) at elevated temperatures, leading to formation of Cu- and S-richer minerals on the crystal surface (Amcoff, 1987, 1988). If the reaction is arrested, while in progress, and the crystal subsequently examined under the microscope a remarkable relationship becomes apparent between the phases. Initially, covellite crystals (CuS) grow outward from the original CuFeS2 grain, while "islands" or clusters of small FeS2 crystals grow around formed "holes" in the CuFeS2 grain. When all sulphur has been consumed, further reaction between CuS and CuFeS2 results in a FeS2 +

Cu5FeS4 assemblage, replacing the original crystal, while CuS is replaced by

Cu5FeS4 only. The interesting point is, that new phases only exhibit physical 24

boundaries towards "older" phases with which they can be in thermodynamic

equilibrium under the experimental conditions, i.e. CuFeS2 + Cu5FeS4 + FeS2 in the interior of the "old" crystal, and CuS + sulphur atmosphere on the surface.

"Intermediate" phases, for ex. Cu5FeSg, also form in an intermediate position. Another point of interest here is that the "old" crystal becomes progressively porous with continuing reaction, showing that the redox process is characterized by outward

diffusion of copper ions essentially, while sulphur from the original CuFeS2 remain

"behind" to form FeS2. a similar process as suggested for natural low temperature

bornite (Cu5FeS4) replacement above.

2.6.2 Low sulphur activity conditions.

If the groundwater is depleted in sulphide species (a small redox capacity), the various Cu-sulphides, which have formed on the canister surface, will start to reequilibrate with respect to each other and to metallic copper. Since dissolved iron species are also expected in the groundwater, it seems likely that this element will also, at least to a degree, be incorporated in the forming sulphides. Assuming an

initial formation of "pure" Cu2S, a subsequent process involving iron will almost certainly be a chemical replacement process; i.e. Fe atoms will replace Cu atoms in the solid matrix. This does not mean a 1:1 process, since Fe has a higher valency (+2 - +3) compared with Cu (+1) in the solid matrix, but the sulphur structural framework will most probably remain more or less unchanged (see bornite replacement, above). Thus, we can expect a process where Cu rapidly diffuses outward through the growing sulphide crystals, while, at the same time, Fe diffuses inward. Since the supply of Cu will be practically unlimited, it seems likely that the forming sulphides will be low in Fe as long as aqueous sulphide species are in abundance. However, if the sulphur supply is depleted while the activity of reactive

Fe remains high in the groundwater Cu2S will continue to incorporate Fe through replacement, and "try to equilibrate" with regard to both metallic copper and the

water phase. Possibly, the net result will be a Cu2S + Cu-j 8S + Cu5FeS4 assemblage, where the Cu-richer phases border the metal. Although thermochemical data on ternary Cu-Fe-S phases are missing, formation of one or several such phases might be inferred from Figure 2. 25

In the present context, it should be mentioned that there are indeed indications that formation of CuFe-sulphides may be kinetically retarded, compared with simple Cu- sulphides. This was suggested by Shea and Helz (1988) who studied the solubility of copper in a reduced sulphidic environment. Besides, they suggested that the solubility of copper decreases strongly in water in contact with CuFe-sulphides compared with water in contact with simple Cu-sulphides. In general, solubilities of metastable phases are higher compared with thermodynamically stable phases (Snoeyink and Jenkins, 1980).

Because of the discussed unknown factors, the present discussion on mineral formation on metallic copper is somewhat speculative. The future histories of path- controlled processes can not be predicted based on theory alone, they can only be inferred from relevant experiments. For example, we would like to draw attention to chlorine, which is expected to be present in the future repository environment.

Unfortunately we have found no data on the behaviour of CI" during sulphidation processes, but we have a strong feeling that it may still be important because of the

chemical similarities with sulphur. F" commonly substitutes for OH" in mineral

structures, by analogy we would expect CI" to substitute for HS" in sulphide structures. Thus, we think that HS-bonds will form, at least to a degree, in all solid sulphides in a hydrous environment, because of simple protonization processes.

There is also a possibility that CI" may substitute directly for S2" in the sulphide

structure, the charge being balanced by Cu1+ vacancy formation, i.e.: S2' + Cu1+ =

CI" + vacancy. Besides, even more possibilities exist since sulphur in some

structures has a negative charge less than 2.

2.7 Rate of sulphidation on metallic copper.

Whatever the exact mechanisms it seems safe to assume that formation of metal sulphides on the surface of metallic copper in a reducing environment will be a rapid

process. In fact, high temperature crystalline modifications of Cu-sulphides, Ag2S (the only known silver sulphide) and members of the ternary Cu-Ag-S system invert so rapidly to low temperature modifications that they are impossible to "freeze". Direct sulphidation experiments also show that copper, like silver, is 26 instantly blackened when put in contact with elemental sulphur, and replacement experiments show that mm-sized crystals of Cu-sulphides and CuFe-sulphides can be inverted into one another within hours below 100°C, if the crystal, is subjected to a suitable concentrated solution (e.g. Schwartz, 1931; Schouten, 1934). These authors did not much consider the mechanisms of the processes, however.

Apparently simple and straightforward sulphidation processes are actually quite complex, when studied in detail, and may involve a number of different steps, where the most sluggish one will determine the rate. Rickard (1975) studied low temperature sulphidation of FeS (approximate formula) to form FeS2 in an aqueous medium, containing sulphide species and S°. (This is considered a typical reducing sedimentary environment). The overall process can be written:

FeS + S° ->FeS2

However, this reaction tells us virtually nothing about the mechanism, since sulphur at low temperature will be present as puckered eight-membered rings

(S8). Thus:

8FeS + S8 ->8FeS2

Such a process is obviously mechanistically impossible as written (a reaction of the 9th order), since it would require simultaneous reaction of all eight sulphur atoms

in a ring. Rickard (1975) performed his rate studies at 40°C, and considered a

process involving dissolution of FeS to form Fe2+(aq) and aqueous sulphide species,

and recombination to form FeS2:

2+ 2 2 + Fe + S5S "(polysulphide) + HS" -> FeS2(pyrite) + S4S " + H

Note that polysulphide ions form through reaction between S° and e.g. S2', and that the aqueous sulphide species here participate in both oxidation and reduction. Thus, considering the mean valency of sulphur: 27

2+ 2 2 + Fe + S5S -(-1/3) + HS"(-2) -> FeS2(-1) + S4S "(-2/5) + H

Rickard(1975) also showed that formation of pyrite at 5°C, in a sedimentary environment containing 1% FeS (particle size around 0.001 mm), 0.01% S° and

PH s=0<0011 would be 2 mgxcm"3x year"1. Sulphidation of Cu-sulphides under similar conditons would probably take place at a much higher rate, since they typically react many magnitudes faster than Fe-sulphides (Barton, 1970).

Besides, Rickard (1975) showed that the rate of "pyritization" was two orders of magnitude higher at 40°C, than at 5°C. Assuming a similar activation energy for sulphidation of metallic copper, this might mean that an initial copper canister temperature of some 80°C (heat produced by decay processes) would be equivalent with a 10 000 times faster process, compared with a 5-10°C milieau.

The examples discussed indicate that the major rate determining factor will not be the sulphidation process itself, but rather the rate of transport of aqueous species to the reaction site on the cansiter surface, i.e. the redox capacity of the groundwater. However, questions like; rate of transport of species through the bentonite barrier, whether the bentonite will survive intact etc., are outside the scope of this report. 28

3 CONCLUSIONS AND RECOMMENDATIONS.

The discussion above has been much focused on path-controlled processes and metastabilities. Several of the examples given are not directly applicable to the problem of mineral formation on metallic copper in a future repository environment. It is necessary, however, to use analogies when studies which have a direct bearing on the problem are missing.

In the geochemical literature it is stressed that low temperature processes seldom result in formation of the chemically most stable product. As indicated by McNeal and Little (1990) and Sato (1992, see Appendix 1), Cu-sulphides may both form and persist, well outside their stability fields. For this reason, and also in view of the chemical complexity of the natural environment, it is quite impossible to predict in detail which chemical reactions will take place and which solid compounds will form. Also, there is some uncertainty regarding a possible formation of Cu-oxides on the canisters. (Perhaps this question has recieved too little attention in the present report.).

However, we think we have indications showing that metallic copper will be subjected to rapid sulphidation /Yreactice aqueous sulphur species are present in the neighbourhood. Iron may also be of importance, as well as other species (CI") which we have not considered. The point is, that the rate of copper "destruction" will depend on the rate with which reactive sulphur species etc. are transported to the canisters. In turn, this depends on the surrounding barriers.

In this report we have not considered factors like; species present in and long-term fate of the bentonite barriers, diffusion rate of reactive species through the bentonite, concentration of species, flow rate of groundwater through neighbouring rocks, microbial activity etc. Here it might be added that microbial life, not dependent on solar energy and photosynthesis, may be widespread at depth in the crust. According to Gold (1992), it may be comparable in mass and volume with all 29 surface life. For the sake of argument, however, we will briefly outline two hypothetic situations, which should cover the whole spectrum of possible futures in a reduced environment: (1) The surrounding rocks do not allow appreciable groundwater flow, and the bentonite barriers act as expected. This means a very small transport of reactive species to the copper canisters. (2) The bentonite- and rock barriers do not act as expected and/or reactive species are released from the bentonite (dissolution of iron sulphide) leading to a high concentraton of reactive sulphur species. (Here we do not consider the possibility that the milieau might accidently become strongly oxidizing.).

(1) After initial reaction with reactive sulphur species etc. in entrapped groundwater, very little will happen (SKBF 83-24). It is a fact that "fresh dense" rocks, free of cracks, can act as very efficient barriers against chemical change. For example, in glacial boulders of quartzite untarnished highly reactive metal sulphides and even metallic copper can be found within centimetres from the surface

(Quartzite is a rock with a very low porosity, made up of quartz; SiO2> crystals. Quartz is resistent against chemical weathering). This means that a barrier, of a few cm of quartzite, has prevented attack from the highly corrosive atmosphere for about 10 000 years.

(2) Since we do not find it likely that forming metal sulphide will form a passivating coating on the canister surface, we believe that the rate of sulphidation will be governed by the rate of transport of aqueous reactants to the canister surface. Besides, since sulphidation experiments suggest a one-way outward diffusion of copper, the canister will simply corrode, and copper in the form of sulphides will reprecipitate on the surface. Nord et al. (1993), who investigated a copper alloy compass ring from SS Kronan, wrecked in the Baltic Sea in 1676, observed that the artefact was covered with a thick crust and that it had no remaining metal core. The

witn crust consisted mainly of spioncopopite (Cu3gS28) traces of covellite (CuS). Diffusion experiments between copper and nickel (Birchenall, 1974) show that metallic copper typically developes pores when one-way outward diffusion takes place. Thus, penetration of the canisters might take place long before all copper has been transformed into sulphide (or been dissolved in the groundwater). Besides, as observed during replacement processes the rate is increased when a loss of volume is 30 involved (Schouten, 1934). Important in this respect are factors like; grain size of the metal, crystal defects, the "quality" of welded joints, texture and initial porosity(?), foreign alloyed elements etc. These factors depend on the mode of manufacturing. Should copper loss be concentrated to a few "high energy points" there is a potential for rapid penetration of the canisters.

In light of a large number of reports by SKB on the expected properties of the surrounding rocks and on the bentonite barrier, it seems likely that formation of metal sulphide on the canisters will be limited. However, considering the uncertainties regarding the chemical reactions involved and the mode of copper loss from the canisters, we would recommend a series of simple sulphidation experiments to be undertaken on metallic copper, in a chemical environment and at a temperature similar to the expected future repository milieau. A better understanding of the mechanisms and rates of mineral formation should hopefully increase our confidence in the use of metallic copper as a technical barrier. Such experiments were outlined in our report SKN 57, 1992. Also, if copper is alloyed, it is important to investigate how this may affect its resistivity against chemical attack. For example, based on numerous observations of low temperature behaviour of Ag-minerals, we have a strong feeling that this element might have an adverse effect on the stability of the canisters.

Parallel with the laboratory experiments we would recommend a series of tests, using nature itself as a laboratory. This means that a number of small copper cylinders, manufactured in exactly the same way, and of identical composition, as the future copper canisters, should be placed in a "future repository site environment", i.e. artificial heating at about 80°C, a bentonite surrounding etc. Such an experiment could be monitored for several decades, and the copper cylinders could be taken out one at a time, after suitable periods of time, and examined for signs of mineral formation and internal corrosion. Even if 10-30 years is a short period of time in the present context, it is not insignificant, because the expected drop in temperature around the canisters will result in a corresponding rapid decrease in reaction rates.

Thus, if copper remains undamaged during the first decades we can be more confident that it will survive as a technical barrier for a long time to come. 31

4 Table 1 Low temperature copper minerals; 1A: Cu-sulphides, 1B: CuFe-sulphides, and 1C: other Cu-phases. Data used here are based on the following references: Bridge,P.J. et al (1978), Bums,P.C. and Hawthome.F.C. (1993), Chao.G.Y. (1981), Craig,J.R.,and Scott.S.D. (1974), Giuseppetti.G. et al (1989), Goble.R.J. (1981), Goble,R.J.,and Robinson.G. (1980), Graedel.T.E. et al (1988), Groat.LA. and Hawthome.F.C. (1987), Jambor.J.L et al (1964),JCPDS (1980), Mumme.W.G. et al (1988), Nickel.E.H. (1973), Pollard.A.M. et al (1989) and (1990), Potter.R.W. (1977), O'Keeffe.M. and Bovin.J.O. (1978), Ridkosil.T. et al (1982), Robertsson.B.E. et al (1967), Rösler.H.J. (1984), Ruotsala.A.P.and Wilson.M.L (1977), Sieber.N.H.W. et al (1987), Van Oosterwyck.G. (1977), Wan.C. et al (1978), Wang.N. (1982), Woods.T.L et al (1986). Further information on mineral stability fields can be obtained from Fig.6 and Fig.7. (Pollard.A.M. et al:1990).

1 A Mineral Composition Structure Th. Stability Remarks C° M i n Max Cu-sulphides

Chalcocite Cu2S Monoclinic 103

Cu2S Hexagonal 103 435

Djurleite Cu1.97S Orthorombic 93

Digenite low CugS5 Cubic 83 stab, by Fe

Cu S 1129 high 9+X 5 Cubic 83

Anilite Cu7S4 Orthorombic 70

Roxbyite Cu1.74-1.82S Monoclinic 65-70 stab, by Fe ass. with sphalerite Geerite Cu1.60S Pseudocubic

Blaubl.covellite Cu1+xS 157

l:Yarrowite Cu1.1+0.1S Hexagonal

ll:Spionkopite Cu1 4S Hexagonal (metastable with re to covellite+anilite)

Covellite CuS Hexagonal 507 32

1 B Mineral Composition Structure Th. Stability Remarks C° Min Max CuFe-sulohides

Digenite (Cu,Fe)9S5 Cubic 83

(Cu,Fe)9S5 Cubic 83 1129

Bornite Cu5FeS4 Tetragonal 228

Cu5FeS4 Cubic metastable forms x-Bornite CuFeS405 Tetragonal 125

Idaite Cu5.5FeS6.5 Hexagonal 501 Nunkundamite Cu5.5FeS6.5

Fukuchilite Cu3FeSg Cubic 200 stable at high P

Chalcopyrite CuFeS2 Tetragonal 557

Cubanite CuFe2S3 Orthorombic 200-210 Interm.sol.sol. Cubic 20-200 960

Talnakhite Cu9Fe8S16 Cubic 186

Interm ph.1 Cu9Fe8S16? ?? 186 230

Mooihoekite Cu9Fe9S16 Tetragonal 167 Interm ph.A CugFegS-15 ?? 167 236 ? Haycockite Cu4Fe5S8 Orthorombic low T Prim. cub. ph. Wide range Cubic? 20 200

?? Cu0.12Fe0.94S1 .00 ?? ????

Valleriite CuFeSo*1.53(Mcj,AI)(OH)o Hex 250

Djerfisherite K3Cu3(FeNi)i -j S14 Cubic ass. with talnakhite, pentlandite and cubanite 33

1C Mineral Composition Structure Th. Stability Remarks C° Min Max Carbonates

Malachite Cu2(CO3)(OH)2 Monoclinic common oi

Azurite Cu3(CO3)2(OH)2 Monoclinic

Nakauriite Cu8(SO4)4(CO3)(OH)6 Orthorombic

Chalconatronite Na2Cu(CO3)2*3H2O Monoclinic corrosion product (synthetic)

Chlorides Nantokite CuCI Cubic Tolbachite CuCl2 Monoclinic unstable in air, hygroscopic

Eriochalcite CuCI2*2H2O Orthorombic ass. with atacamite

Botallackite Cu2(OH)3CI Monoclinic

Paratacamite Cu2(OH)3CI Trigonal 40

Atacamite Cu2(OH)3CI Orthorombic

Claringbullite Cu8(OH)14Cl2*H2O

Calumetite Cu(OH,CI)2*2H2O Orthorombic

Anthonyite Cu(OH,CI)2*3H2O Monoclinic

Connellite Cu37CI8(SO4)2(OH)62*8H2O Hexagonal

Buttgenbachite (SO4,NO3)—- Hexagonal

Mitscherlichite Tetragonal synthesized at room temp. (synthetic) 34

1C Mineral Composition Structure Th. Stability Remarks - C° Min Max Oxides and Hvdroxides

Cuprite Cu2O Cubic Paramelaconite CU4O3 Tetragonal ass. with tenorite Te no rite QD Monoclinic

Delafossite CuFeO2 Trigonal

Cuprospinel CuFe2O4 Tetragonal 700-800 stable with hematite

Cubic 700-800 stab, by Mg at room temp.

Spertiniite Cu(OH)2

Gerhardtite Cu2(OH)3NO3 Orthorombic 140 ass. with malachite, cuprit

Phosphates

Cu-pyrophosph .a-Cu2P2O7 Monoclinic 70-100

b-CupP2O7 70-100

Sampleite CU3(PO4)2*HpO Orthorombic ass. with atacamite, cor-

rosion product on artifacts

Libethenite CUOPOA(OH) Orthorombic ass. with malachite

Cornetite Cu3PO4(OH)3 Orthorombic

Reichenbachite Cu5(PO4)2(OH)4 Monoclinic ass. with ps.malachite

Pseudomalachite Cus(PO4)2(OH)4 Monoclinic

(*H2O)

Tagilite CU2(PO4)OH*H2O? Monoclinic ass. with libethenite.and

ps.malachite.

Hentschelite CuFe2(PO4)2(OH)2 Monoclinic

Chalcosiderite CuFe6(PO4)4(OH)s Triclinic (Cornwall)

*4H2O Likasite Orthorombic ass. with gerhardtite 35

1C Mineral Composition Structure Th. Stability Remarks C° Min Max Silicates

Chrysocolla Orthorombic

Dioptase CuSiO2(OH)2 Trigonal

Shattuckite Cu5(SiO3)4(OH)2 Orthorombic Plancheite Orthorombic ass. with dioptase, malachite,cuprite

(Bisbeeite) =plancheite+chrysocolla (Katangite) =microcryst. variety of plancheite(?) Kinoite Monoclinic in skarn and in basalt ass. with native copper

Papagoite CaCuAISi2O6(OH)3 Monoclinic ass. with ajoite Litidionite Triclinic (Vesuvius) Ajoite Triclinic 425

(OH)6*3H2O 36

1C Mineral Composition Structure Th. Stability Remarks C° Min Max Sulphates

Chalcocyanite CUSO4 Orthorombic (=Hydrocyanite)

Poitevinite CuSO4*H2O Monoclinic theoretical formula

Bonattite CuSO4*3H2O Monoclinic

Chalcanthite CuSO4*5H2O Triclinic

Boothite CuSO4*7H2O Monoclinic

Antlerite Cu3SO4(OH)4 Orthorombic Brochantite Cu4SO4(OH)6 Monoclinic ass. with cuprite, also common on artifacts

Posnjakite Cu4SO4(OH)6*2H2O Monoclinic

Langite Cu4S04(0H)6*2H2O Orthorombic

Wroewolfeite Cu4SO4(OH)6*2H2O Monoclinic

Dolerophanite Monoclinic converts to antlerite(?) in

water at room temp.

Guildite CuFe(SO4)2(OH) Monoclinic

•4H2O

Cuprocopiapite CuFe4(SO4)6(OH)2 Triclinic

*20H2O

Devillite Cu4Ca(SO4)2(OH)6 Monoclinic

•3H2O

Natrochaicite NaCu2(SO4)2OH*H2 0 Monoclinic Krohnkite • Monoclinic

Cyanochroite(s)K2Cu(SO4)2*6H2O Monoclinic (s)=synthetic sample Leightonite l<2Ca2Cu(SO4)4*2H2O Triclinic (Pseudoorthorombic)

Chlorothionite K2CUSO4CI2 Orthorombic water soluble 37

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Appendix 1.

6 CONSTRAINTS ON THE USE OF Eh-pH STABILITY DIAGRAMS TO PREDICT FORMATION OF SULPHIDE MINERALS.

6.1r General.

Much of the discussion below is based on Sato (1992). When discussing sulphide mineral relations in a natural water environment at low temperature, diagrammatic presentations of phase stability fields in terms of two or more intensive variables (intensive v. = thermodynamic v. not related to mass), such as temperature, pressure, and chemical potentials (related to activities), are of great help. Most common is the use of Eh-pH diagrams. However, when these equilibrium diagrams are confronted by the real world, large differences frequently become apparent. For example, as discussed at length by Sato (1992); large discrepancies are observed, when phase relations in the Cu-Fe-S-0 system at room temperature, constructed by Mckinstry (1959) from field- and microscopic observations of supergene sulphide - and oxide ore, are compared with phase relations in the same system, as calculated from thermochemical data (reproduced after Garrels and Christ, 1965).

As discussed above, this can partly be explained by the fact that many mineral phases which have experienced changes at low temperature in nature, owe their present appearence to metastable processes. However, it has been observed, as well, that redox processes between aqueous species, where several electrons are exchanged, and/or where a considerable structural reorganization of the molecule takes place, are very sluggish at low temperatures. As a result of this, direct potentiometric measurements often give results that differ strongly from values calculated from thermodynamic equilibria. (Perhaps most well-known in the present context is

2 2 reaction between SO4 " and S ", the kinetics of which have been investigated by

Ohmoto and Lasaga, 1981, who used exchange of 34S - 32S between the oxidized and

reduced species, and who reported insignificant rates below 200°C in a non- bacterial environment). Below, we will give a few examples of redox disequilibria, which may have a bearing on the present problem. 46

According to Sato (1992), the reaction between oxygen and water is the most striking example of redox disequilibrium that persists for geological time in the Eh range of subareal waters, especially ocean water. This very important redox process defines the upper Eh boundary of the natural environment. Under standard conditions the equilibrum between oxygen and water is given by:

+ O2 + 4H + 4e" = 2H2O

Eh = 1.229 + 0.0148logao - 0.0591pH (A1)

At the atmosphere - water boundary (PQ _ 0.21 atm.) this becomes: 2~

Eh = 1.22 - 0.059pH (A2)

Measured potentials of various aerated waters are all about 0.4 volt lower than, predicted from this equation (Garrels and Christ, 1965). Thus, another control mechanism must be in effect. According to Sato (1960, cited in Sato, 1992) the observed persistent non-equilibrium is due to the high activation energy needed to break the strong bonds in molecular oxygen to form "atomic" oxygen. A number of authors, including Sato (1992), have considered an alternative reaction:

+ O2 + 2H + 2e" = H2O2

Eh = 0.682 - 0.0591ph + 0.0295logao /aH Q

The details of hydrogen peroxide formation in natural waters, which are complex, have been investigated by Millero et al (1987). It has also been shown to be of importance for the Cu1+/Cu2+ redox couple, i.e. for the persistence of Cu1+(aq) in seawater (Moffet and Zika, 1983).

However, according to Sato (1992), reduction of hydrogen peroxide takes place through: 47

+ H2O2 + 2H + 2e" = 2H2O

Eh = 1.772 - 0.0591pH + 0.0295logaH 0 (A4)

As seen in these reactions, the Eh at which H2O2 is reduced to H2O is much higher than its oxidation Eh under identical conditions, which means that H2O2 is unstable with respect to decomposition into H2O + O2. This process can be catalyzed by redox couples with intermediate E° values, for example Fe3+/Fe2+, with E° = 0.77 volt,

2+ 3+ where Fe is oxidized and Fe reduced by H2O2. leading to little net change in the

Fe3+/Fe2+ ratio.

If Eh of a solution containing ions of multivalent elements, like Fe and Mn, is lower than in equation A3, both O2 and H2O2 can act as oxidizers, resulting in a rapid increase in Eh. When Eh exceeds that of equation A3, however, O2 ceases to be an effective oxidizer, because H2O2 starts to decompose into O2 and H2O as soon as it forms. This means that H2O2 no longer contributes to the net oxidation, while equation A2 is blocked by a high energy barrier. According to Sato (1992) the empirical upper limit of Eh of aerated waters, containing elements like Fe and Mn corresponds to a ag /aH o ra*'° °' "^ 'n ecluation A3, which gives:

Eh = 0.86 - 0.059pH (A5)

Highly oxidized minerals can form, however, in spite of the discussed constraint on

Eh by the O2/H2O2 process. This is so, because soiid oxides containing mixed valency elements (Fe, Mn) can disproportionate if the chemical environment is suitable. According to Sato (1960), simple acidification can bring about disproportionate of intermediate Mn-oxides to Mn2+ ions plus "higher" oxides. For example:

+ 2+ Mn2O3 + 2H = MnO2 + Mn + H2O

2+ where Eh is finally controlled by the MnO2/Mn redox couple: 48

+ 2+ MnO2 + 4H + 2e" = Mn + 2H2O

Eh = 1.228 - 0.1182pH - 0.0295logaMn2+ (A6)

At variance with oxides, metal sulphides behave in a complex manner when subjected to changing redox conditions. This is so, because both metal atoms and sulphur may participate in electron exchange processes. This property leads to stability relations, which, at first glance, may seem contradictory. For example, when metallic copper is oxidized in an environment where reactive sulphur species and water are present, copper sulphides form in order of increasing oxidation; Cu° ->

Cu2S -> -> CuS -4 Cu2S -» Cu° (see for example, Figure 5, page 60-61 in SKN 59, 1992). For simplicity we have ignored all copper sulphides of composition intermediary between Cu2S (chalcocite) and CuS (covellite). If we assume that copper is univalent in all copper sulphides (personal communications with Rolf Berger, Institute of Chemistry, University of Uppsala) the average valency on sulphur changes according to: Cu° -> Cu2S(-2) -> -4 CUS(-1) ->...-> Cu2S(-2)

2 -> Cu°. This is explained by the fact that we assume SO4 "(aq) to be in perfect equilibrium with the copper sulphides and with the present aqueous sulphide species. This is not the case at low temperature.

The present example is instructive in showing that the chemical milieau, in terms of Eh, can not be inferred from valencies of metal atoms and sulphur in metal sulphides. Moreover, as already mentioned above, metal sulphide relations observed in nature often differ significantly from what is expected from equilibrium considerations. This is the topic under the next heading. 49

6.2 Persistency field Eh-pH relations.

The discussion below is to a large degree based on Sato (1992) and references therein. Persistency field Eh-pH relations differ from thermodynamic equilibrium relations in that observed metastability fields are included. Thus, persistency relations should resemble more closely what actually takes place in nature, and, hence, be of value in predicting reaction paths.

Direct potentiometric measurements are employed to deliniate the metastability fields of semiconducting metal sulphides. The solid sulphide electrode is connected via an external circuit to an inert metal electrode (Pt, Au), and both electrodes are immersed in the solution. When Eh of the solution is higher than that of the solid sulphide, the latter will release electrons which are transported via the external circuit to the inert metal electrode, where reduction takes place; e.g. Fe^+ + e" -»

Fe2+. When Eh of the solution is lower than that of the metal sulphide, the latter will be reduced and recieve electrons via the external circuit etc.

It is important to note that in the "natural system" the solution is in most cases regarded as having a much larger redox capacity than the metal sulphide, i.e. it is continually replenished from outside. (If we assume that the technical barriers function as expected, the situation will be the opposite as regards the copper canisters, thus, the solution Eh will change. This is not discussed here).

Now, the actual reaction is identified by a comparison of the observed phase- boundary electrode potential (=potential where the solid phase is oxidized and reduced, respectively) with computed potentials of various possible electrode reactions. Thus, it is observed that oxidation takes place via release of electrons and metal ions, resulting in formation of a sulphide with a lower metal to sulphur ratio. If such a phase does not exist, elemental sulphur may form.

For example, for a Cu2S electrode in an acid, slightly oxidizing solution, the reversible phase-boundary reaction observed is in good agreement with the following model: 50

2+ CuS + Cu + 2e" = Cu2S

Eh = 0.578 + 0.0295logaCu2+ (A7) and for a CuS electrode:

S° + Cu2+ + 2e" = CuS

Eh = 0.591 + 0.0295logaCu2+ (A8)

Electrode reactions of copper sulphides with compositions intermediate between

Cu2S and CuS should fall between these values. In a near-neutral to weakly basic solution, acu2+ kecomes a function of pH because of hydroxide or oxide formation.

Hence for the Cu2S and CuS oxidation:

+ CuS + CuO + 2H + 2e" = Cu2S + H2O

Eh = 0.753 - 0.0591pH (A9) and:

+ S° + CuO + 2H + 2e" = CuS + H2O

Eh = 0.812 - 0.0591pH (A10)

On the other hand, reduction of copper sulphide is observed to proceed via a release of sulphur as sulphide ions into the solution and formation of a sulphide with a higher metal to sulphur ratio, according to the stability Eh-pH diagram (finally metallic copper will form).

2 Thus, for reduction of CuS in the S ', HS", and H2S predominance fields, respectively. 51

2 2CuS + 2e' = Cu2S + S "

Eh = -0.538 - 0.0295logas2- (A11)

+ - 2CuS + H + 2e = Cu2S + HS"

Eh = -0.124 - 0.0205logaHS- - 0.0295p H (A12)

+ 2CuS + 2H + 2e" = Cu2S + H2S

Eh = 0.083 - 0.0591pH - 0.0295logaH s (A13)

(again ignoring intermediary copper sulphides).

To summarize; observed oxidation of metal sulphides leading to a release of metal ions is path-controlled and results in metastable equilibria. On the other hand, reduction reactions, where sulphur is released as sulphide ions, are in concordance with the stability Eh-pH relations.

The situation becomes more complex, when we try to define the electrode potential of a metal sulphide in the Eh-pH region where it can be in equilibrium with the solution. Sato (1966) derived a theoretical expression for a simple binary sulphide,

+ with the stoichiometric formula; M2S;, which dissociates in the solution into M) and S2" ions, as:

2 2 Eh = E° + 2.303(RT/4jF)log[(aMj+) (as)J/(as2-)J(aM) ] (A14)

(see Appendix in SKN 57, 1992, for definitions)

E° can be computed using: 52

+ o 2 E° = (1/4jF)[2AG°f(MJ ) + jAG f(S ")] (A15) where AG°f(MJ+) is the Gibbs free energy of formation of MJ+ at 25°C and 1 atm., etc. The relation between Gibbs free energy and the equilibrium constant was discussed in our report SKN 57, 1992. The activity of the metal ion; a^j+, and the sulphide ion; ag2-, are related to each other through the solubility product (Ks):

2 2 M,2Sj = 2MJ+ + jS ", Ks = (aMj+) (as2-)J (A16) where:

+ 2 logKs = -[2AG°f(Mi ) + jAG°f(S -) - AG°f(M2Sj)]/2.303j) R T (A17)

The activities of metal and sulphur in the metal sulphide; aM and as, are related to each other by means of the solid state activity product (Ka):

2 M2Sj = 2M+jS Ka = (aM) (as)J (A18) where:

logKa =AG°f(M2Sj)/2.303RT (A19)

It is important to note that the components in the solid; M and S, are not equivalent with the the ionic species; MJ+ and S2", in the solution. Also note that the solid state solubility product (which is not often used, since it is difficult to calculate activities in solid compounds) is derived from the Gibbs-Duhem relation. Thus, for a binary compound:

n1d)i1 + n2dn2 = 0 (A20)

where \L-\ and n2 are the chemical potentials and n-j and n2 are the number of moles, 53 of components 1 and 2, respectively, in a binary solid solution (Klotz and Rosenberg, 1986).

The chemical potential can be related to the activity:

|ij = n,° + RTInaj -» d^ = RTdlnaj (A21)

Where \i-\ is the chemical potential of component i in the solid solution, n°j is the chemical potential of component i in the pure phase i, and a\ is the activity of component i in the solid solution. substitution yields:

n^lna-i + n2dlna2 = 0 (A22)

Integration yields: n^ Jdlna-j = -n

Ina-j = -n2lna2 + C or:

ni n2 In[(a1) (a2) ] = C (A23) or:

n l n2 c (a1) ' (a2) = e = Ka (solid state solubility product) (A24)

However, for CuS; logaCu varies from -6.62 to -8.60 and logas from -1.98 to 0. (It is assumed that the most sulphur-rich CuS is in equilibrium with pure sulphur

at 25°C, 1 atm.). For Cu2S, logaCu varies from 0 to -6.62 and logas from -15.22 54 to -1.98 (It is assumed that the most copper-rich Cu2S is in equilibrium with pure metallic copper). Again, intermediary copper sulphides have been ignored.

In one important aspect, the oxidation trend observed in potentiometric measurements contrasts with the trend expected from •thermodynamic stability

2 considerations. Thus, the latter trend indicates that formation of SO4 " directly from a metal sulphide, gives the lowest computed oxidation potential among

2 alternative possible oxidation paths, at a given pH. This means that "SO4 ' forming" reactions define the lowest possible upper boundary of sulphide stability fields, regardless of whether these reactions really take place at the solid surface. Also, these reactions predict that the oxidation potential of a sulphide should become more positive as the SO^2" activity in the solution increases.

In reality, however, the potential of a metal sulphide is not directly affected by aSO 2"- ((unlesun'ess sosoli''dd rnetametall ssulphate is precipitated, in which case the activity of the metal in solution is affected).

The reason for the observed behaviour of semiconducting metal sulphides during oxidation is that a direct formation of SO^2' is inhibited by a high activation energy, while release of metal as an aqueous ion is a much simpler process. It may also be noted that thiosulphate ions; S2O32", act as a metastable intermediary at low temperatures. (Incidently, if we consider the metastable redox process; S° -»

S2O32", which has been indicated by thiobacterial oxidation; Zajic, 1969, cited in Sato, 1992, elemental sulphur will also have a metastable extension towards higher Eh values, although it can not be directly measured potentiometrically, since sulphur is an insulator).

When a persistency field Eh-pH diagram for a copper sulphide is compared with a corresponding stability field it is obvious that the solid phase may persist under higher Eh conditions than expected from purely thermodynamic considerations

(Figure 8). Besides, CuS can persist at higher Eh than Cu2S, which is also at variance with what we would expect from equilibrium considerations. Within the persistency field, the metal sulphide can survive for long periods of time at room 55 temperature, at least as regards redox reactions. However, processes that do not require destruction of the solid sulphur structural framework are possible, such as cation exchange and other replacement processes including redox processes involving sulphur. 56

7 FIGURE CAPTIONS.

Figure 1: Phase relations in the central parts of the "dry" Cu-Fe-S system at decreasing temperatures. Note the increase in complexity (compare with Figure 2), and the decrease in extents of solid solutions, with decrease in temperature. The tie- lines connect phases which are believed to be in thermodynamic equilibrium under the specified conditions (temperature specified, total pressure equivalent with the sum of all partial vapor pressures, for a specific phase assemblage). In the diagrams we can distinguish one-phase, two-phase, and three-phase assemblages, depending on bulk composition. Abbrevations: dgss: digenite solid solution, bnss: bornite solid solution, cv: covellite (CuS), id: idaite (Cu5FeS6), cp: chalcopyrite (CuFeS2), iss: intermediate solid solution, poss: pyrrhotite solid solution, py: pyrite (FeS2). Based on a number of authors, refe:red to in the text.

Figure 2: Tentative room temperature relations in the Cu-Fe-S system. Two possibilities are shown: (a) tie-line digenite - pyrite (Cu-| 8S - FeS2), (b) tie-

line idaite - chalcopyrite (Cu5FeS6 - CuFeS2). Note that other possibilities exist as well. The diagrams are simplified along the binary joins Cu-S and Fe-S.

Abbrevations: cc: chalcocite (Cu2S), dg: digenite (Cu-| 8S), an: anilite (Cu-j 75S),

cv: covellite (CuS), id: idaite (Cu5FeSg), bn: bornite (CU5F6S4), cp: chalcopyrite

(CuFeS2), t: talnakhite (CugFe8S16), m: moohekite (Cu9FegS16), h: haycookite

(CugFe-iQS-jg), cb: cubanite (CuFe2S3), tr: troilite (FeS), po: monoclinic low- temperature pyrrhotite (Fe7S8), py: pyrite (FeS2). Based on a number of authors, referred to in the text.

Figure 3: The chalcopyrite (CuFeS2) structure, which can be considered a typical metal sulphide structure where the metal atoms in a regular way occupy tetrahedral intersticies between cubic close packed sulphur atoms. The packing arrangement is identical with that of sphalerite (ZnS). Instead of the simple rows of Zn atoms in ZnS the cation positions are occupied alternatively by Cu and Fe atoms. The result is

that the unit cell of CuFeS2 corresponds to two unit cells of ZnS stacked one on top of the other. Note that although the symmetry is tetragonal the basic packing arrangement is still face-centered cubic. 57

At 557°C (P = 0) the CuFeS2 structure transforms to a "true" sphalerite structure since the Cu and Fe atoms are now statistically distributed among the cation positions.

Figure 4: Examples of solid state replacement as seen through the microscope.

(a) Massive replacement where Cu5FeS4 has formed as a result of solid state diffusion between Cu2S and CuFeS2. The simplified overall process can be written:

CuFeS2 + 2Cu2S -> Cu5FeS4 However, to illustrate that the process takes place via cation diffusion:

+ - 2+ 2CuFeS2 + 3Cu + e -> Cu5FeS4 + Fe , and

2+ + 4Cu2S + Fe -» Cu5FeS4 + 3Cu + e"

Note that a chemical gradient develops through the growing Cu5FeS4 zone; it is Cu- rich at the Cu2S boundary and Cu-poor at the CuFeS2 boundary, resulting in a change in colour from bluish (Cu-rich) to brownish (Cu-poor). After Amcoff (1988). (b) Oriented replacement or replacement exsolution with lamellae of the replacing phase(s) growing in the host grain. In the text two examples of oxidation are discussed, both of which result in a similar texture, i.e:

+ (i) Cu5FeS4 -> CuFeS2 + 2Cu1 8S + 0.4Cu + 0.4 e", and

2+ (ii) 3Fe3O4 -> 4Fe2O3 + Fe + 2 e"

In (i) duplex lamellae of CuFeS2 + Cu-| 8S form along (111)-planes in a Cu5FeS4 host crystal, and in (ii) lamellae of Fe2O3 form along (111)-planes in a Fe3O4 host crystal (the martitization process).

+ 2+ Note that Cu ions precipitate as CuS and Fe ions as Fe2O3, respectively, outside the grains which are being replaced! This is due to the "inertness" of sulphur and oxygen, respectively, compared with the cations as regards solid state diffusion.

Figure 5: Hypothetic sulphidation of metallic copper in an environment containing aqueous sulphide species, e.g. HS", H2S, polysulphide ions. (a) Initial conditions.

(b) Polysulphide ions may act as electron acceptors and release S2' at the metal surface. 58

+ (c) Chalcocite; Cu2S, is formed initially. Outward diffusion of Cu through the metal

2 and the growing metal sulphide is followed by combination with S " to form Cu2S.

2 2 For the overall redox process e.g: 2Cu° + S5S " -» Cu2S + S4S ". (d) With continuing reaction S-richer Cu-sulphides are formed, depending on the total concentration of dissolved sulphide species. A mineralogical steady state gradient may be set up, with Cu2S in contact with the metal and S-rich Cu-sulphide in contact with the water-rich environment. Pores develop and grow in the metal, because of preferential copper depletion along defects etc. (e) If the solution is depleted in reactive species the sulphide crystal layer will equilibrate with respect to both metallic copper and solution chemistry. If Fe(aq) species are present CuFe-sulphides may form through replacement of "simple" Cu- sulphides.

Figure 6: Eh-pH diagram for some "oxidized" copper minerals

3 1 <75 (a) with aSO42-(aq)=10- and aCL-(aq)=1°'

2 ac =10 2 and 1 75 (b) with aso4 -( l) ' aCL-(aq)=1Cr - Note that in (b) only minerals whose stability fields have been changed are named.

Figure 7: Stability field diagram for the copper(ll)halide minerals, (logacr as a function of pH)

(a) in the abscence of CO2

3 5 (b) CO2(g) with a partial pressure of 1(T - (Normal atmospheric conditions).

36 (c) CO2(g) with a partial pressure of 1(H-

In all three diagrams the sulphate activity is equal to 10"3 Note that in (b) and (c) only minerals whose stability fields have been changed are

named.

Figure 8: Persistency fields for CuS (a) and Cu2S (b). Cu-sulphides of intermediate compositions have been neglected for simplicity. Note that reduction

reactions (CuS -> Cu2S; Cu2S -> Cu°) correspond with equilibrium processes (the lower persistency boundaries in the diagrams), while the oxidation reactions are path-controlled, and correspond with a release of copper and formation of more S- 59

2+ rich phases (Cu2S -» CuS; CuS ->S° + Cu ). See the text for details. Thus, a significant overvoitage is needed to "destroy" Cu-sulphides (like other metal sulphides!), once they have formed.

Total concentration of dissolved sulphur species: 10"1 molar. Note that this is only of importance for the lower persistency boundaries! Simplified after Sato (1992). 600°C

three-phase

two-phase one-phase

bn,

/ Cu 400°C py Fe

ss

250°C py

SS

Cu Fe

Fig. 1 py Room temperature (tie-line, dg-py) cv

Cu Fe

Room temperature (tie-line, id-cp)

Cu

Fig. 2 CuFeS2 (low T) (Cu, Fe)2S2 (high T)

O's Fe Ocu Statistical Cu, Fe distribution

Fig. 3 CuFeS2 Cu2S Initial conditions

50 jim

a) massive replacement

oriented b) replacement

Fig. 4 a) \ b) \ \ \ \ \ environment \ B \ with reduced CD S-species, e.g; _ \ CD H2S, HS", \ Q. Q. polysulphidic ions \ \ and Fe2+ etc. \ 8 \ \ \ \ \ \ \

Fig. 5 Paratacamite Connellite 0.4 ^Ci^OHfeCI Cu37CI8(SO4)2(OH)62

x8H2O O iCuC I " 12 0.3 -oI •E CD Z

ri t a) o Cuprite ^^ c iS 0.2 CD . (Cu2O) ^\

0.1 - Cu

0 i i 6 8 pH

CD o B x o'- •H Q. % nJ •«• o Antlerit e 1 § CO Cu37CI8(SO4)2(OH)62 ^<4 x8H2O

Cu O b)

LU . Cu2O ^\

Cu

i 5 6 7 8 9 pH

Fig. 6 Paratacamite

Cu2(OH)3CI CuO, 2-

-1 14 pH

Fig. 6(a) 0.5-

0-

Eh

-0.5-

-1

pH

Fig. 8(b) STATENS KÄRNKRAFTINSPEKTION Swedish Nuclear Power Inspectorate

' "' • • I

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